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General Chemistry

First Semester Final Exam Study Guide

59 multiple choice questions

1. Define product.

2. Define reactant.

3. Label the products and reactants in the following chemical reactions:

H2 (g) + O2 (g) ( H2O (l)

Copper sulfate(aq) + iron(s) ( iron sulfate(aq) + copper(s)

H2O + CO2 + light ( C6H12O6 + O2

4. Define element and give three examples.

5. Define compound and give three examples of compounds.

6. Explain the law of conservation of mass and give a scenario that supports this law.

7. Complete the chart.

|Subatomic Particle |Charge |Size |Location in Atom |

| | | | |

| | | | |

| | | | |

8. Which subatomic particle makes up most of an atom’s mass?

9. Most of an atoms volume is really what?

10. Define isotope and give an example.

11. Why is the atomic mass of a given element usually NOT listed on the periodic table as a whole number?

12. What particle does the atomic number listed on the periodic table represent, and where does this particle reside?

13. Write the equation that relates the mass # to the number of protons and neutrons in the nucleus of an atom.

14. Complete the chart.

|Name |Symbol |Protons |Neutrons |Electrons |Mass Number |Atomic Number |

|Potassium | | | | | | |

| |Pb | | | | |82 |

|Arsenic | | | |33 | | |

| | |47 |63 | | | |

| | | | | |260 |101 |

15. Which elements from problem 36 were isotopes?

16. List the four electron orbitals AND the total number of electrons each orbital can hold.

17. Identify the element with 4 energy levels, 5 valence electrons and is a metalloid.

18. Draw the orbital diagrams for the following atoms:

a. Carbon

b. Oxygen

c. Aluminum

19. Write the complete (unabbreviated) electron configuration for the following atoms:

a. Helium

b. Oxygen

c. Aluminum

d. Neon

20. Write the abbreviated (noble gas) electron configuration for the following atoms:

a. Silver

b. Copper

c. Tungsten

d. Titanium

21. Define ground state.

22. Define excited state.

23. How do atoms jump from ground state to excited state and back? (Remember our Flame Test Lab!)

24. Which sub-atomic particle changes energy states?

25. Who is considered the “father” of the modern periodic table?

26. Define period.

27. Define group.

28. Define electronegativity.

29. Define atomic radius.

30. Define ionization energy.

31. Label the following on the attached periodic table (page 13):

1) Highest electron affinity

2) Highest electronegativity

3) Largest atomic radius.

4) Highest ionization energy.

5) Highest reactivity.

6) Alkali metal

7) Alkaline earth metals

8) Halogens

9) Noble gases

10) Actinide

11) Lanthinide

12) Metals

13) Nonmetals

14) Metalloids

15) Solids

16) Liquids

17) Gases

18) Valence electrons for groups 1, 2, 13-18

32. Fill in the Table Below

|Type of Element |Conductivity |Reacts with Acid |State of Matter |Malleable |Brittle |

| |YES OR NO |YES OR NO | |YES OR NO |YES OR NO |

| | | | | | |

|Metal | | | | | |

| | | | | | |

|Nonmetal | | | | | |

| | | | | | |

|Metalloid | | | | | |

33. Label the subscripts in the following chemical formula.

K2SO4

34. What information does a subscript reveal and how would changing a subscript in a chemical formula change the chemical it represents?

35. Define molar mass.

36. Define molecular formula.

37. Define empirical formula.

38. Calculate the molar mass of the following:

a) Au

b) H2O

c) Mg(MnO4)2

d) CH3CH2COOH

e) Fe2(SO4)3

39. What is the empirical formula for C6H12 ?

40. What is the empirical formula for Hg2Cl2 ?

41. What is the percent composition of each element for the following compounds:

a) potassium cyanide, KCN

b) butane, C4H10

c) sulfuric acid, H2SO4

42. Find the empirical formula of a compound that is 63.52 % iron and 36.48 % sulfur.

43. Find the empirical formula of a compound that contains 32.38 % sodium, 22.65% sulfur, and 44.99% oxygen.

44. Calculate the empirical formula of a compound containing 1.0 g K, 0.70 g Cr, and 0.82 g of O.

45. Determine the molecular formula of a compound with an empirical formula of CH2 and a formula mass of 42.08 g/mol.

46. If you have a mole of pennies, how many do you have?

47. Convert 3.5 moles of nickel into grams.

48. What is the mass of 5 moles of C6H12O6, glucose?

49. How many grams are in 3 moles of HCl, hydrochloric acid?

50. Convert 100 grams of CaCO3 (chalk) into moles?

51. How many moles are in 1 gram sample of gold?

52. How is an ion different then an atom?

53. Define cation and give an example. Do cations form from metals or nonmetals? Explain.

54. Define anion and give an example. Do anions form from metals or nonmetals? Explain.

55. Define ionic bond and give an example.

56. Name the following ionic compounds.

a) Al2(SO4)3

b) PbF4

c) KBr2

d) Ca(OH)2

e) Mg(NO2)2

57. Write the formula for the following ionic compounds.

a) Magnesium phosphide

b) Calcium cyanide

c) Tin (IV) chloride

d) Zinc nitrate

e) Calcium fluoride

58. Define covalent bond.

59. Name the following covalent (in other words molecular) compounds.

a) As3P5

b) IF7

c) NO2

d) SrH2

e) SiCl4

60. Write the formula for the following covalent (in other words molecular) compounds.

a) dinitrogen pentoxide

b) carbon monoxide

c) trisulfur difluoride

d) triarsenic pentanitride

e) dihydrogen monoxide

61. Compare and contrast ionic and covalent bonds (include the properties).

62. Define polar covalent bond and give an example.

63. Define nonpolar covalent bond and give an example.

64. How many electrons are shared in a single covalent bond?

65. Define octet rule. Which family on the period table has a complete octet and thus does not react to form compounds?

66. Define valence electron.

67. How many valence electrons do the following atoms have:

a) Magnesium

b) Bromine

c) Calcium

d) Phosphorous

68. The electron configuration of oxygen is 1s2 2s2 2p4. How many more electrons does oxygen need to satisfy the octet rule?

69. When an atom loses or gains electrons to become an ion, how many electrons will be in the outermost energy level?

70. How do you balance a chemical reaction?

71. Balance the equation below:

____ Al + ____ CuSO4 ( ____ Al2SO4 + _____ Cu

72. In a chemical formula, what do the subscripts tell you?

73. Balance the equations below:

____ H2 + _____ O2 ( _____ H2O

____ (NH4)NO2 ( _____ N2 + _____ H2O

74. What does a balanced chemical equation tell you?

75. Identify the type of reaction (synthesis, double displacement, etc) below:

Ca + FeO ( CaO + Fe

Mg + S ( MgS

2 NaI + MgCl2 (2 NaCl + MgI2

76. Define the following:

a) Synthesis-

b) Single displacement-

c) Decomposition-

d) Double displacement-

77. What type of reaction is it, if your products are CO2 + H2O?

78. Define activity series and explain how it works.

79. Predict the products of the following reactions using the activity series in your textbook:

a) Cu(SO4) + Fe (

b) CaF2 + H2 (

80. In the following reaction, 2 NaI + MgCl2 (2 NaCl + MgI2, what is the mole ratio of MgCl2 to NaCl?

81. Write the calculation for a grams-grams conversion (mass-mass).

82. If you have 22. 9 g of Ni and 112 g of AgNO3, which reactant is limiting? Which is excess?

2 AgNO3 + Ni => 2 Ag + Ni(NO3)2

83. Using the reaction and your calculations from number 82, what mass of Ni(NO3)2 would be produced?

84. Write the calculation for moles-gram conversion (mole-mass).

85. What happens to a reaction when the limiting reactant is used up?

86. Write the calculation for a grams-mole conversion (mass-mole).

87. Write the calculation for a mole-mole conversion.

88. If you have 5.94 g of Cu and 23.23 g of HNO3, which reactant is limiting? Which is excess?

4 HNO3 + Cu => Cu(NO3)2 + 2 H2O + 2 NO2

89. Using the reaction and your calculations from number 88, what mass of NO2 would be produced?

90. CuCl2 + 2 Na(NO3) ( Cu(NO3)2 + 2 NaCl

If 15 grams of copper (II) chloride react with 20 grams of sodium nitrate, how many grams of sodium chloride can be formed?

91. Using the balanced equation in #90, how many moles of CuCl2 are reacted to form 3 moles of NaCl?

92. How many moles of aluminum will be produced from 30,000.0 g of Al2O3 in the following reaction?

2 Al2O3 ( 4 Al + 3 O2

93. How many moles of Fe2O3 will react with 99 grams of Al?

Fe2O3 + 2 Al ( 2 Fe + Al2O3

94. What mass of P will be needed to produce 3.25 moles of P4O10?

4 P + 5 O2 ( P4O10

95. What mass of O2 is produced when 1.840 moles of H2O2 decomposes?

2 H2O2 ( 2 H2O + 1 O2

96. Write the formula for calculating percentage yield.

97. For the reaction: 2 Fe(PO4) + 3 Na2(SO4) → Fe2(SO4)3 + 2 Na3(PO4), if I perform the reaction with 25 grams of iron (III) phosphate, how many grams of iron (III) sulfate can I make?

98. Using the reaction in #97, if 18.5 grams of iron (III) sulfate are actually produced, what is my percent yield?

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